If you're seeing this message, it means we're having trouble loading external resources on our website.

If you're behind a web filter, please make sure that the domains *.kastatic.org and *.kasandbox.org are unblocked.

Gas mixtures and partial pressures

For a mixture of ideal gases, the total pressure exerted by the mixture equals the sum of the pressures that each gas would exert on its own. This observation, known as Dalton's law of partial pressures, can be written as follows:  P (total) = P ₁ + P ₂ + P ₃ + ... where P ₁, P ₂, and P ₃ are the partial pressures of the different gases in the mixture, and P (total) is the total pressure of the mixture.. Created by Sal Khan.

Want to join the conversation?

• At why is the total pressure 2.5atm, where did that number come from and how did you calculate it?
• The number 2.5atm was just made up as a starting value for the question, and wasn't calculated from anything. The part that was calculated was "what is the partial pressure of the oxygen in the container, since the pressure changed from 2.0atm to 2.5 atm when we added oxygen in with the nitrogen?"

Hope that helps.
• Hello. When you add molecules of oxygen gas to the nitrogen gas, why does the number of moles (n) stay the same?
• Yes, when we added O2 molecules, we did change the number of moles in the container. However, in the video, She is saying that the number of moles of N2 didn't change since we only added some molecules of O2. Pressure is dependent on 3 factors. (i.e. T, V and n) Since we didn't change any of these factors for N2, n of N2 stays the same.
• Does it not make a difference wether the particles are colliding with each other and so the temperature increases which leads to the change in pressure?
• It is an assumption made that the collisions between the particles are completely elastic and that is the reason the avg. energy of particles remains constant, hence there is no change in temperature too.
Hope that helps.
• Air is made up of different types of gases and all these gases considered to be ideal.
They all are in same space so the factor V is same for all.
They are at the same temp, and being similar in mass, volume since they are "ideal" they all would have same average KE. So T is also same for all.
The only thing however is different is n.
So doesn't that mean just knowing n of each gas can give us their Partial pressure?
• Why is R the same? Don't different gases have different gas constants? Or is R the same for simplicity, to make the problem easier to answer?
• The gas constant, R, is the same for all gases.

Anytime you here that is something is proportional to something else, like x is proportional to y, then another way to say that is that x = ky where k is some constant.
In physics & chemistry we figure out that we have proportional relationships: but then we need to figure out how exactly they are proportional, which we figure out through experimentation.

Remember that the gas law was discovered using 3 observations of 3 people:
Boyle's Law --> V ∝ 1 / P,
Charles's Law --> V ∝ T,
Avogadro's law --> V ∝ n

Then, we wanted to find some way to combine these 3 proportional relationships, and so we did that using the gas law.
V = nRT/P

Note that we are really just combining the 3 said relationships, and we put the constant there to define the proportionality. Also it is often seen as PV = nRT.

Hope this helps,
- Convenient Colleague
• there was initially 4 nitrogen molecules exerting a P of 2 atm then we added 4 oxygen molecules but the total P wasn't 4 atm.
This means O and N have different P, and since V and n is same, then they must be different in terms of T.
But won't the temperature of both N and O be same after a while? If that is so then after a while N and O won't have the same Partial pressure we calculated it has just after adding them together?
• You're taking the diagrams in the video too literally. They are meant to be symbolic and they don't accurate represent the amounts of each gas.

Although the diagrams suggest that the number of molecules (or moles) of nitrogen equal those of oxygen, this cannot be the case for an ideal gas, given the pressures referred to in the video. For the total pressure to be 2.5 atm after adding the oxygen, then there must be four times the amount of nitrogen than there is oxygen.
• If i am given the total pressure, how can i get the partial pressure of the individual gases that make the gas mixture since the gases apply different pressure on the container
• You can use the mole fraction of each gas in the mixture to find their partial pressures if you're given total pressure and the moles of gas present. Sal shows this at .
• I know according to the Ideal Gas Equation, pV=nRT, the removal or addition of gases from a mixture does not affect its partial pressure. But if we see the formula for partial pressure which is mole fraction multiplied by the total pressure. And mole fraction is the number of moles of a particular molecule divided by the total number of moles.

By removing or adding any gases, wouldn't the total number of moles change? Hence causing the mole fraction and partial pressure to change as well..
• I think you have a conceptual misunderstanding. Adding or decreasing gasses does affect the partial pressure. The exception is when adding an INERT gas which doesn't change anything.
• A sample of unknown gas is collected above a closed rigid container of water
at 80o C, the total pressure above the container is 125 kPa. What is the
pressure due to the unknown gas?
• I usually don't like answering obvious homework questions since it's the easy way out, but I don't like hanging people out to dry either so I guess you're in luck.

You have to use Dalton's law of partial pressures which states that the total pressure of mixture of gases is the combination of the component gases' partial pressures added together. Partial pressures being the pressure exerted by a single gas compound.

Here you have two gases, an unknown gas and water vapor (gaseous water) in a rigid container. It's rigid so the gases won't expand the container and the pressure remains constant. The total pressure is 125 kPa and is composed of the partial pressures of the unknown gas and water vapor. So the formula looks like this so far using Dalton's law: Ptotal = 125 kPa = Punknown +Pwater. So you have two variables, the Punknown and Pwater, and they want to eventually solve for the Punknown which means they gave you a way to determine the Pwater in the problem. We can determine the Pwater from the temperature which gives us the vapor pressure of water. The vapor pressure is the pressure of the evaporated gaseous water above liquid at a certain temperature in a closed container. So the vapor pressure of water for a temperature is a constant you can look up in a table anywhere. Using Wikipedia is acceptable here which tells us the vapor pressure for water at 80 degrees Celsius is 47.3730 kPa. So 47.3730 kPa is your Pwater for the equation from before. Therefore, 125 kPa = Punknown + 47.3730 kPa. Using algerbra to solve for Punknown yields us 77.627 kPa. Acconting for sig figs, the final answer is 78 kPa.

Hope that helps.